Rate of Reactions | O-Level Chemistry

Last updated: 9 October 2026

Diagnose → Learn → Master: Every study-note section includes 15 free MCQs — start with 5 Diagnostic MCQs to identify gaps in your understanding, then complete 10 Mastery MCQs near the end to check what you have learned.

Some reactions seem almost instantaneous, while others take minutes, days or even years. Magnesium can burn rapidly in oxygen, while iron rusts much more slowly. Both are chemical reactions, so what makes one reaction happen much faster than another?

To answer that question, we need to connect two views of the same reaction. In the laboratory, we measure how quickly something changes. At particle level, we explain that change by considering how particles move, collide and react. Once those two views are connected, factors such as concentration, temperature, pressure, surface area and catalysts become much easier to understand.

Before we build that explanation, try the short diagnostic quiz. It is not meant to test everything in the chapter yet. Its purpose is to reveal which ideas you already understand and which ones deserve closer attention.

Diagnostic quiz · 5 questions

How Well Do You Know Rate of Reactions?

Try these questions before reading the chapter. They check the basic ideas needed to understand reaction rates.

Identify what you already understand.
Spot common misconceptions early.
Read the chapter with a clearer purpose.

Which statement best describes a faster chemical reaction?

It always produces a greater final amount of product.

A faster reaction does not necessarily produce more product overall. The final amount depends on how much limiting reactant is present.

A reactant is used up, or a product is formed, more quickly.

Correct. Reaction rate describes how quickly the amount of a reactant decreases or the amount of a product increases.

It must release more heat.

Reaction rate and energy change are different ideas. A reaction can be fast or slow whether it is exothermic or endothermic.

It must occur at a higher temperature.

Higher temperature can increase rate, but reactions can differ in rate for many other reasons, including concentration, surface area and catalysts.

Reaction rate is about how quickly chemical change occurs, not how much product is eventually obtained.

A reaction produces a gas. Which apparatus directly measures the volume of gas produced?

Electronic balance

A balance measures mass. It can be used for a mass-loss rate experiment, but it does not directly measure gas volume.

Burette

A burette is designed to deliver accurately measured volumes of liquid.

Gas syringe

Correct. Gas entering the syringe pushes the plunger outwards, allowing its volume to be read directly.

Thermometer

A thermometer measures temperature rather than gas volume.

A gas syringe is especially useful in rate experiments because the gas volume can be recorded at regular time intervals.

Which statement about collisions between reactant particles is correct?

Particles must collide effectively for a reaction to occur.

Correct. Reactant particles must collide, and the collision must meet the conditions needed for reaction.

Every collision between reactant particles produces products.

Many collisions do not result in reaction because they may not have sufficient energy or a suitable collision arrangement.

Particles can react without coming into contact.

Collision theory requires reacting particles to collide before their atoms can rearrange.

Only collisions between solids can cause reactions.

Collisions are important in reactions involving gases, solutions and solid surfaces as well.

Collision theory explains reaction rate by considering how often reactant particles undergo effective collisions.

Why can a more concentrated solution react faster than a dilute solution at the same temperature?

The particles become larger.

Concentration does not change the size of the reactant particles.

Each particle gains more kinetic energy.

At the same temperature, increasing concentration does not give each particle more kinetic energy.

The activation energy becomes lower.

Increasing concentration does not lower the activation energy.

There are more reactant particles per unit volume.

Correct. More particles in the same volume leads to more frequent collisions and hence more effective collisions per second.

Concentration changes how many reactant particles occupy a given volume. It does not make each particle move faster.

What is the main way in which a catalyst increases the rate of a reaction?

It increases the temperature of the reactants.

A catalyst can work without increasing the temperature.

It provides a reaction pathway with lower activation energy.

Correct. The lower activation energy allows a greater proportion of collisions to result in reaction.

It increases the amount of limiting reactant.

Adding a catalyst does not add more limiting reactant.

It permanently becomes part of the products.

A catalyst remains chemically unchanged overall at the end of the reaction.

A catalyst changes the pathway of the reaction rather than supplying extra energy to the particles.

What Does Rate of Reaction Mean?

The rate of reaction is a measure of how quickly a reactant is used up or how quickly a product is formed.

To measure a rate, we therefore need two pieces of information:

a measurable change during the reaction; and
the time over which that change occurs.

Average rate = change in measured quantity ÷ time taken

The unit of rate depends on what is being measured. If gas volume is measured, a possible unit is cm³ s⁻¹. If mass loss is measured, the unit could be g s⁻¹.

A measurement becomes a rate only when time is included.

Saying that 60 cm³ of gas was produced tells us the amount of gas formed. Saying that 60 cm³ was produced in 30 s tells us how quickly it formed.

Worked example: calculating an average rate

Calcium carbonate reacts with dilute hydrochloric acid:

CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)

Suppose 48 cm³ of carbon dioxide is collected in 30 s.

Average rate = 48 ÷ 30 = 1.6 cm³ s⁻¹

This is the average rate during those 30 seconds. The reaction does not necessarily produce exactly 1.6 cm³ every second. In most reactions, the rate changes as the reaction proceeds.

How Can Reaction Rate Be Measured?

There is no single apparatus that measures the rate of every reaction. Instead, choose a quantity that changes clearly as the particular reaction proceeds.

Measuring the volume of gas produced

If a reaction produces a gas, the gas can be collected in a gas syringe. Its volume is then recorded at regular time intervals.

For the reaction between calcium carbonate and hydrochloric acid, for example, carbon dioxide may be measured every 10 seconds.

The apparatus must be airtight.

If carbon dioxide escapes through a loose bung or leaking connection, the gas syringe records less gas than the reaction actually produced.

Measuring the decrease in mass

A reaction that produces an escaping gas can also be followed by placing the reaction flask on an electronic balance.

As gas leaves the flask, the total mass decreases. Recording the mass at regular intervals shows how quickly the gaseous product is being formed.

Loose cotton wool can be placed in the neck of the flask to reduce the loss of liquid droplets through splashing while still allowing the gas to escape.

Gas syringe or mass loss?

A gas syringe directly measures the volume of gas produced.

A mass-loss experiment measures the decrease in mass as a gaseous product escapes.

Both can be used to follow a reaction over time, but they measure different quantities.

Timing the formation of a precipitate

Some reactions produce an insoluble solid called a precipitate. As more precipitate forms, the mixture may become increasingly cloudy.

One example is the reaction between sodium thiosulfate and dilute hydrochloric acid. Solid sulfur forms:

Na₂S₂O₃(aq) + 2HCl(aq) → 2NaCl(aq) + SO₂(g) + S(s) + H₂O(l)

The flask can be placed over a black cross. The time taken for the cross to disappear from view is recorded.

A faster reaction produces enough sulfur to hide the cross in a shorter time. When the same endpoint is used in every experiment:

Relative rate = 1 ÷ time taken

If the cross disappears in 20 s:

Relative rate = 1 ÷ 20 = 0.050 s⁻¹

This is a relative rate measurement.

The experiment does not measure the exact mass of sulfur produced. It compares how long different reactions take to reach the same chosen level of cloudiness.

Choosing a suitable measurement

What changes during the reaction?	Suitable method	Important consideration
A gas is produced	Measure gas volume using a gas syringe	Make sure the apparatus is airtight
A gaseous product escapes	Measure the decrease in mass using a balance	Prevent splashing while allowing gas to escape
A precipitate makes the mixture cloudy	Measure the time taken to obscure a mark	Use the same endpoint and viewing conditions

Collision Theory: Why Do Reactions Happen?

Measuring a reaction tells us how fast it occurs. Collision theory helps us explain why the rate changes.

Reactant particles must first collide with one another. However, simply bumping into one another is not always enough.

Effective collisions

A collision that results in a chemical reaction is called an effective collision.

For a collision to be effective:

the particles must collide with sufficient energy; and
the parts of the particles that need to react must meet in a suitable arrangement.

The second condition is sometimes called suitable orientation. Molecules are not simply featureless balls: particular atoms or bonds may need to meet one another during the collision. If the wrong parts of the particles meet, the particles can collide with enough energy and still separate without reacting.

Comparison of ineffective and effective collisions showing insufficient collision energy, unsuitable collision arrangement and a successful collision with sufficient energy and suitable arrangement — A collision produces a reaction only when the particles have sufficient energy and the reacting parts meet in a suitable arrangement.

Activation energy

The activation energy, Eₐ, is the minimum energy that colliding particles must possess for a reaction to occur.

If the particles collide with less than this energy, the collision is ineffective and the particles separate without forming products.

Activation energy is an energy barrier.

It is not the amount of energy released by the reaction. Even an exothermic reaction may need an initial energy input before it can begin.

For example, methane releases energy when it burns, but it does not normally ignite at room temperature without a flame or spark:

CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l)

What determines the reaction rate?

A reaction becomes faster when there are more effective collisions per second.

This can happen because:

particles collide more frequently;
a greater proportion of collisions have sufficient energy; or
both effects occur together.

Factors That Affect the Rate of Reaction

The different factors affecting reaction rate may seem like separate facts to memorise. They become much easier when each one is linked back to collision theory.

1. Concentration of reacting solutions

Increasing the concentration of a reactant means that there are more reactant particles per unit volume.

The particles are therefore more likely to meet one another:

there are more reactant particles in the same volume;
collisions occur more frequently;
there are more effective collisions per second; and
the reaction rate increases.

Higher concentration does not make the particles move faster.

If temperature remains constant, their kinetic-energy distribution does not increase simply because the solution is more concentrated. The important change is the number of reactant particles per unit volume.

Example: magnesium reacting with hydrochloric acid

Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)

The same piece of magnesium reacts faster in a more concentrated hydrochloric acid solution because there are more reacting particles per unit volume available to collide with the magnesium surface.

Example: combustion in oxygen

Increasing the concentration of oxygen can also increase the rate of combustion. Magnesium, for example, burns more vigorously in pure oxygen than in air.

2Mg(s) + O₂(g) → 2MgO(s)

In pure oxygen, there is a much higher concentration of O₂ molecules around the magnesium. Collisions between oxygen molecules and the magnesium surface are therefore more frequent.

2. Pressure of gaseous reactants

Pressure is especially important when the reacting substances are gases.

Increasing the pressure by compressing a gas places the particles into a smaller volume.

there are more gas particles per unit volume;
the particles collide more frequently;
there are more effective collisions per second; and
the rate increases.

Concentration and pressure use the same collision idea

Increasing concentration: more dissolved reactant particles are present per unit volume.

Increasing gas pressure: gaseous reactant particles are pushed into a smaller volume.

Both changes increase collision frequency.

3. Surface area of a solid reactant

When a solid reacts with a solution or gas, reaction occurs at the surface of the solid.

Breaking the same amount of solid into smaller pieces increases its total surface area.

More solid particles are then exposed to the other reactant at the same time. Collisions at the solid surface therefore occur more frequently.

Powder compared with large pieces

Powdered calcium carbonate reacts more rapidly with hydrochloric acid than the same mass of large calcium carbonate chips.

The powder has a greater total surface area, so more calcium carbonate particles are exposed for collision.

Comparison of the same mass of calcium carbonate as large chips and smaller particles showing that smaller particles have a greater total surface area and undergo more frequent collisions with acid — Breaking the same mass of solid into smaller pieces increases the total exposed surface area, allowing more collisions with the acid each second.

Smaller pieces do not contain more substance if the mass is unchanged.

The important difference is that more of the solid is exposed at the surface. The reaction can therefore occur at more places simultaneously.

If equal masses of calcium carbonate react completely with excess acid, the powder reacts faster but does not produce more carbon dioxide overall. Both samples contain the same amount of calcium carbonate.

4. Temperature

Increasing temperature has two effects on the particles.

Particles collide more frequently

At a higher temperature, particles have greater kinetic energy and move faster. They therefore collide more frequently.

A greater proportion of particles have sufficient energy

This is the particularly important part of the explanation.

At a higher temperature, a greater proportion of particles possess energy equal to or greater than the activation energy.

A greater proportion of collisions can therefore be effective.

For temperature, give both collision-theory effects.

The particles collide more frequently, and a greater proportion of them have sufficient energy to overcome the activation-energy barrier.

Why “the particles move faster” is incomplete

Faster-moving particles do collide more frequently, so that statement is not wrong. However, it leaves out the important activation-energy effect.

A stronger explanation states that the proportion of collisions with sufficient energy to react also increases.

Summary of the main rate factors

Factor increased	Particle-level change	Why the rate increases
Concentration	More reactant particles per unit volume	More frequent collisions
Pressure of gaseous reactants	Gas particles occupy a smaller volume	More frequent collisions
Surface area of a solid	More solid particles are exposed	More collisions can occur at the solid surface
Temperature	Particles move faster and more have sufficient energy	More frequent collisions and a greater proportion of effective collisions

Catalysts and Activation Energy

A catalyst is a substance that increases the rate of a chemical reaction while remaining chemically unchanged at the end of the reaction.

A catalyst works by providing an alternative reaction pathway with a lower activation energy.

At the same temperature:

the catalyst lowers the activation-energy barrier;
a greater proportion of collisions have enough energy to react;
there are more effective collisions per second; and
the reaction becomes faster.

A catalyst provides an alternative pathway with a lower activation energy. The reactant and product energy levels remain unchanged.

A catalyst does not give the particles more kinetic energy.

It makes it easier for the existing collisions to result in reaction by providing a pathway that requires less activation energy.

Example: decomposition of hydrogen peroxide

Hydrogen peroxide slowly decomposes into water and oxygen:

2H₂O₂(aq) → 2H₂O(l) + O₂(g)

Manganese(IV) oxide, MnO₂, increases the rate of this reaction and remains chemically unchanged overall.

What a catalyst does not change

It does not increase the amount of limiting reactant.
It does not increase the final amount of product obtainable from the same starting amounts.
It does not permanently become part of the products.
It does not work by increasing the temperature of the particles.

Interpreting Reaction-Rate Graphs

Rate experiments often produce graphs showing how a measurable quantity changes with time.

Common examples include:

volume of gas produced against time;
mass of product formed against time;
mass of the reaction mixture against time; and
amount of reactant remaining against time.

The gradient tells us the rate

The gradient of a graph tells us how quickly the measured quantity is changing.

Gradient = change in vertical quantity ÷ change in time

A steeper graph represents a faster rate of change.

For a product–time graph, the graph slopes upwards as product forms. For a mass-loss or reactant–time graph, the graph may slope downwards. In both cases, a steeper slope means a faster reaction.

Why is the reaction usually fastest at the beginning?

At the beginning, the reactant concentrations are at their highest. Reactant particles therefore collide most frequently.

As reactants are used up, their concentrations decrease. Effective collisions occur less frequently, so the reaction becomes slower.

Why does the graph eventually become horizontal?

A horizontal section has zero gradient. The measured quantity is no longer changing, so the reaction rate has fallen to zero.

This normally happens because a reactant has been completely used up.

The particles have not stopped moving.

A horizontal graph means there is no further overall chemical reaction because one of the required reactants is no longer available.

Steepness and final height tell us different things

Suppose equal masses of calcium carbonate react completely with excess hydrochloric acid. One experiment uses powder and the other uses large chips.

The powder gives a steeper graph because it reacts faster. However, both graphs eventually reach the same final carbon dioxide volume because both experiments contain the same amount of calcium carbonate.

Reaction-rate graphs showing that a steeper gradient represents a faster reaction while the final plateau height represents the total amount of product formed — The gradient shows the rate of reaction, while the final plateau shows the total amount of product formed. A faster reaction does not necessarily produce more product.

Do not confuse rate with amount

Gradient or steepness: tells you about reaction rate.

Final height: tells you about the total amount of product formed.

A reaction can therefore be faster without producing more product overall.

When can the final height change?

The final amount of product changes when the amount of the limiting reactant changes.

For example, doubling the amount of calcium carbonate while keeping hydrochloric acid in excess can produce more carbon dioxide. The final plateau would then be higher.

When comparing graphs, ask two separate questions.

First ask which curve is steeper and therefore faster. Then ask whether the amount of limiting reactant has changed, because that determines the final amount of product.

Planning a Fair Rate Experiment

A rate experiment is useful only when the effect of one factor can be separated from the effects of all the others.

Suppose a student investigates how hydrochloric acid concentration affects its rate of reaction with calcium carbonate.

Identify the variables

Variable	Example in the investigation
Independent variable	Concentration of hydrochloric acid
Dependent variable	Volume of CO₂ produced at different times
Control variables	Mass and surface area of CaCO₃, acid volume, temperature and apparatus

A suitable procedure

Measure a fixed volume of hydrochloric acid of known concentration.
Measure a fixed mass of similarly sized calcium carbonate chips.
Place the calcium carbonate into a reaction flask.
Connect the flask to a gas syringe using an airtight bung and delivery tube.
Add the acid and start the stopwatch immediately.
Record the volume of carbon dioxide at regular time intervals.
Repeat using different concentrations of hydrochloric acid.
Keep all relevant control variables constant.
Repeat trials where appropriate and calculate a mean.

Two ways to compare the reactions

You can compare:

the volume of gas produced after the same time; or
the time taken to produce the same volume of gas.

Same time or same amount?

At the same time, the reaction that has produced more gas is faster.

For the same gas volume, the reaction that takes less time is faster.

Common Mistakes in Rate-of-Reaction Questions

Common mistake	Why it is wrong	Better explanation
“There are more collisions.”	It does not clearly state that the collisions occur more frequently.	There are more frequent collisions, giving more effective collisions per second.
“Higher concentration makes particles move faster.”	Particle kinetic energy depends mainly on temperature.	Higher concentration means more reactant particles per unit volume, so collisions occur more frequently.
“Powder contains more particles than the same mass of lumps.”	Equal masses contain the same amount of substance.	Powder has a greater total surface area, so more particles are exposed for collision.
“Higher temperature only causes more collisions.”	This misses the important activation-energy effect.	Collisions are more frequent and a greater proportion of particles have sufficient energy to react.
“A catalyst gives particles more energy.”	A catalyst does not increase particle kinetic energy.	A catalyst provides an alternative pathway with lower activation energy.
“A faster reaction produces more product.”	Rate and final amount of product are different ideas.	A faster reaction has a greater rate. The final product amount depends on the amount of limiting reactant.
“The graph is horizontal because the particles stop moving.”	Particles still have kinetic energy.	The reaction has stopped because a required reactant has been completely used up.

Rate of Reactions Checklist

Measuring rate

Can I identify a measurable change during the reaction?
Can I calculate change ÷ time?
Can I choose between gas volume, mass loss and precipitation methods?
Can I state a suitable unit for the rate?

Collision theory

Can I explain what an effective collision is?
Can I define activation energy?
Can I distinguish collision frequency from collision success?
Can I link a faster rate to more effective collisions per second?

Rate factors

Can I explain concentration using particles per unit volume?
Can I explain pressure for gaseous reactants?
Can I explain surface area using exposed solid particles?
Can I give both effects of increasing temperature?
Can I explain how a catalyst lowers activation energy?

Graphs and experiments

Can I use gradient to compare reaction rates?
Can I explain why reactions slow down?
Can I distinguish final product amount from reaction rate?
Can I identify independent, dependent and control variables?

Mastery quiz · 10 questions

Rate of Reactions Mastery Quiz

Now apply the chapter to calculations, experimental results and collision-theory explanations.

Apply ideas rather than simply recall definitions.
Practise interpreting experiments and reaction graphs.
Check whether your collision-theory explanations are exam-ready.

A reaction produces 72 cm³ of gas in 40 s. What is its average rate of gas production?

Average rate = change in gas volume ÷ time

0.56 cm³ s⁻¹

This does not result from dividing 72 cm³ by 40 s.

1.2 cm³ s⁻¹

This is too small. Use the measured change in gas volume divided by the full time taken.

1.8 cm³ s⁻¹

Correct. 72 ÷ 40 = 1.8 cm³ s⁻¹.

2.9 cm³ s⁻¹

This is larger than the value obtained from 72 ÷ 40.

For an average rate, divide the measured change by the time taken: 72 ÷ 40 = 1.8 cm³ s⁻¹.

Two experiments produce carbon dioxide. Curve X is steeper than curve Y, but both curves level off at the same final gas volume. What can be concluded?

X is faster, but both experiments eventually produce the same amount of gas.

Correct. The steeper curve shows a greater rate, while the same plateau shows the same final amount of product.

X is faster and must contain more limiting reactant.

If X contained more limiting reactant, it would normally be able to form a greater final amount of product, giving a higher plateau.

Y is faster because it takes longer to reach the plateau.

Taking longer indicates a slower reaction, not a faster one.

The reactions must have identical rates because the final volumes are equal.

Equal final amounts do not mean equal rates. Rate is shown by the gradient.

Gradient tells you about rate; final height tells you about total product. These must be interpreted separately.

A student investigates how hydrochloric acid concentration affects its rate of reaction with calcium carbonate chips. Which factor is most important to keep constant so that surface area does not affect the comparison?

The colour of the hydrochloric acid

Hydrochloric acid is colourless, and colour is not a rate factor in this experiment.

The brand of stopwatch

A suitable stopwatch can measure time regardless of brand. This does not control the solid surface area.

The size of the reaction flask only

Using consistent apparatus is sensible, but this does not directly control the exposed surface of the calcium carbonate.

The mass and size of the calcium carbonate chips

Correct. Keeping the chips similarly sized prevents changes in surface area from becoming another variable.

The experiment should change only acid concentration. The amount and surface area of calcium carbonate must therefore be controlled.

Equal masses of powdered calcium carbonate and large calcium carbonate chips each react completely with excess hydrochloric acid. Which result is expected?

CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)

The chips react faster and produce more CO₂.

Large chips have a smaller total surface area, so they react more slowly.

The powder reacts faster, but both produce the same final amount of CO₂.

Correct. Powder has a larger surface area, but equal masses contain the same amount of CaCO₃.

The powder reacts faster and produces more CO₂.

Greater surface area changes the rate, not the number of moles present.

Both react at the same rate because their masses are equal.

Equal mass does not mean equal surface area. The powder exposes much more solid to the acid.

Surface area changes how quickly the reaction occurs. The amount of CaCO₃ determines how much CO₂ can ultimately form.

Which explanation best accounts for the increase in reaction rate when temperature is increased?

The particles become larger and therefore collide more often.

Heating does not make the particles themselves larger.

The activation energy automatically becomes lower.

Increasing temperature does not lower the activation energy of the pathway.

Particles collide more frequently and a greater proportion have sufficient energy to react.

Correct. Both increased collision frequency and the activation-energy effect contribute to the faster reaction.

Every collision becomes effective at the higher temperature.

Even at higher temperature, many collisions may still be ineffective.

A complete temperature explanation includes two effects: more frequent collisions and a greater proportion of particles with energy equal to or greater than Eₐ.

The pressure of gaseous reactants is increased by compressing them at constant temperature. Why does the reaction become faster?

The particles occupy a smaller volume and collide more frequently.

Correct. Compression increases the number of gaseous particles per unit volume and hence the collision frequency.

The particles gain kinetic energy because pressure has increased.

The question states that temperature remains constant, so the increase in rate is not explained by greater kinetic energy.

The activation energy becomes zero.

Increasing pressure does not remove the activation-energy requirement.

Each gaseous particle becomes heavier.

Compression does not change the mass of an individual particle.

Pressure affects gaseous reaction rate mainly through collision frequency.

A catalyst is added to a reaction without changing the temperature. Which statement is correct?

All reactant particles now have greater kinetic energy.

The temperature has not changed, so the catalyst does not increase the particles' kinetic energy.

The catalyst increases the maximum amount of product.

A catalyst changes how quickly products form, not the amount obtainable from the same quantities of reactants.

The catalyst increases the activation energy so collisions are more energetic.

A catalyst does the opposite: it provides a pathway with lower activation energy.

A greater proportion of collisions can overcome the lower activation-energy barrier.

Correct. More collisions can now result in reaction at the same temperature.

A catalyst changes the reaction pathway, allowing a greater proportion of collisions to be effective.

In a disappearing-cross experiment, Experiment P takes 20 s while Experiment Q takes 40 s to reach the same endpoint. How do their relative rates compare?

Relative rate = 1 ÷ time

P has half the relative rate of Q.

The shorter time corresponds to a greater relative rate, not a smaller one.

P has twice the relative rate of Q.

Correct. 1/20 is twice 1/40.

P and Q have equal relative rates.

Their times to reach the same endpoint are different, so their relative rates are different.

Q has four times the relative rate of P.

Q takes longer, so it is slower rather than faster.

For a fixed endpoint, rate is inversely related to time. Halving the time corresponds to doubling the relative rate.

In a mass-loss rate experiment, why is a loose plug of cotton wool sometimes placed in the neck of the reaction flask?

To stop all gas from leaving the flask

Gas must escape for the mass of the apparatus to decrease.

To increase the concentration of the reactants

Cotton wool does not increase reactant concentration.

To reduce loss of liquid droplets while still allowing gas to escape

Correct. This makes the measured mass decrease mainly because gaseous product leaves rather than because liquid splashes out.

To act as a catalyst for the reaction

The cotton wool is not being used as a catalyst.

The cotton wool improves the experiment only if it remains loose enough for gas to escape.

A graph of gas volume against time becomes horizontal after 90 s. What is the best interpretation?

No more gas is being formed because a required reactant has been used up.

Correct. The horizontal section has zero gradient, so there is no further measurable reaction.

All particles in the flask have stopped moving.

Particles continue to move. The graph is horizontal because no further overall reaction is occurring.

The activation energy has increased to infinity.

The plateau is explained by reactant depletion, not an infinite activation energy.

The reaction is continuing at a constant maximum rate.

A horizontal graph has zero gradient, so the measured rate is zero rather than maximum.

A horizontal product–time graph indicates that product is no longer being formed. A required reactant has been completely used up.

Frequently asked questions

Questions Students Often Ask About Reaction Rates

These questions bring together the practical measurements, collision theory and reaction graphs that often cause confusion when they are studied separately.

Is reaction rate the same as the time taken for a reaction?

No. They are related, but they are not the same quantity. If two experiments are compared using the same fixed endpoint, the reaction that reaches that endpoint in the shorter time is faster. This is why relative rate is sometimes expressed as 1 ÷ time.

Why can we use 1 ÷ time in a disappearing-cross experiment?

Every trial is stopped when the same chosen level of cloudiness hides the cross. The amount of visible change is therefore treated as fixed. A shorter time means a faster reaction, so 1 ÷ time gives a convenient relative measure of rate.

Does increasing concentration give the particles more energy?

No. If temperature remains constant, increasing concentration does not make each particle move faster. Instead, there are more reactant particles per unit volume, so collisions occur more frequently.

Why is temperature different from concentration when explaining reaction rate?

Increasing concentration mainly increases collision frequency by placing more reactant particles in each unit volume. Increasing temperature changes particle kinetic energy. The particles collide more frequently, and a greater proportion of them have sufficient energy to overcome the activation-energy barrier.

Does using powder instead of lumps change the amount of product formed?

Not if the same amount of solid reacts completely and the other reactant is in excess. Powder has a greater surface area and therefore reacts faster, but the same number of moles of limiting reactant produces the same final amount of product.

Why does a catalyst not increase the final amount of product?

A catalyst changes how quickly the reaction proceeds by lowering the activation energy. It does not add more limiting reactant. Therefore, the same starting amounts can produce the same maximum amount of product, but they reach it more quickly.

Why is a reaction usually fastest at the beginning?

The reactant concentrations are greatest at the beginning. Reactant particles therefore collide most frequently. As reactants are used up, their concentrations fall and the reaction slows.

Can a downward-sloping graph still represent reaction rate?

Yes. A mass-loss graph or a graph showing reactant remaining may slope downwards. A steeper downward slope still represents a faster change. What matters is the magnitude of the gradient, not whether the line slopes upwards or downwards.

Why can two reaction graphs have the same final height but different shapes?

They can contain the same amount of limiting reactant but react at different rates. For example, powdered calcium carbonate and the same mass of large chips can produce the same final amount of CO₂, but the powder reaches that amount more quickly because it has a greater surface area.

What should I write when an exam asks me to explain why a reaction is faster?

State what changes at particle level, then connect that change to the number of effective collisions per second. For example, for greater concentration: there are more reactant particles per unit volume, collisions occur more frequently, and therefore there are more effective collisions per second.

Reaction rate becomes much easier when every explanation returns to the same particle-level question: what causes more effective collisions to occur each second? Concentration, pressure and surface area mainly change how frequently particles can collide. Temperature also changes how many particles have sufficient energy, while a catalyst lowers the activation-energy barrier. The experiments and graphs we observe in the laboratory are simply the visible consequences of those particle collisions.